Volume and temperature relationship of a gas – Charles' law - Pass My ExamsPass My Exams Volume and temperature relationship of a gas – Charles' law Volume and temperature relationship of a gas Charles' law The relationship between the volume and temperature of a gas was first put forward by the French scientist Jacques-Alexandre-César Charles at around 1787 and is known as Charles’ Law. Charles’ law states: "For a fixed mass of gas, at a constant pressure, the volume (V) is directly proportional to the absolute temperature (T)." Volume α Temperature The animation below gives and explanation of Charles' law: A sealed cylinder with no leaks contains a fixed mass. In order to keep the gas pressure constant the piston is allowed to move freely so that the internal pressure created by the gas particles can equal the constant external pressure. If the internal pressure increases the piston will move up to allow the pressure to equalise. The above set up is used to investigate the relationship between temperature and volume for a gas. Heat energy is applied to the cylinder and the temperature of the gas increases. The average velocity of the gas particles increases resulting in an increase in the rate of collisions and the average force per collision.

This produces an increase in pressure inside the cylinder, the cylinder pressure becomes greater than the external pressure and the piston moves up increasing the volume. By plotting the recorded values of volume (V) against temperature (T) a straight line is produced. We can see from the values that the gas expands uniformly with temperature. We can extrapolate the straight line and see the relationship between cooling the gas and the volume. Further extrapolation gives the temperature at which the volume of gas would become zero. This temperature is at -273°C and is called the absolute zero of temperature. Converting the recorded temperatures into the Kelvin scale and plotting the volume (V) against the absolute temperature (T) gives a straight line which when extrapolated passes through the origin. This shows the volume of the gas is directly proportional to the absolute temperature of the gas. Doubling the temperature will double the volume. The gradient of the slope is the constant in Charles’ Law. Charles’ Law Example: Using the example of the sealed cylinder above, the volume of gas at the start is recorded as 30 cm3 with a temperature of 30°C. The cylinder is heated further till the thermometer records 60°C. What is the volume of gas?

Solution: Gas Laws Gas Laws One of the most amazing things about gases is that, despite wide differences in chemical properties, all the gases more or less obey the gas laws. The gas laws deal with how gases behave with respect to pressure, volume, temperature, and amount. Pressure Gases are the only state of matter that can be compressed very tightly or expanded to fill a very large space. Pressure is force per unit area, calculated by dividing the force by the area on which the force acts. The earth's gravity acts on air molecules to create a force, that of the air pushing on the earth. This is called atmospheric pressure. The units of pressure that are used are pascal (Pa), standard atmosphere (atm), and torr. 1 atm is the average pressure at sea level. It is normally used as a standard unit of pressure. The SI unit though, is the pascal. 101,325 pascals equals 1 atm. For laboratory work the atmosphere is very large. A more convient unit is the torr. 760 torr equals 1 atm. A torr is the same unit as the mmHg (millimeter of mercury).

It is the pressure that is needed to raise a tube of mercury 1 millimeter. The Gas Laws: Pressure Volume Temperature Relationships Boyle's Law: The Pressure-Volume Law Boyle's law or the pressure-volume law states that the volume of a given amount of gas held at constant temperature varies inversely with the applied pressure when the temperature and mass are constant. Another way to describing it is saying that their products are constant. PV = C When pressure goes up, volume goes down. When volume goes up, pressure goes down. From the equation above, this can be derived: P1V1 = P2V2 = P3V3 etc. This equation states that the product of the initial volume and pressure is equal to the product of the volume and pressure after a change in one of them under constant temperature. For example, if the initial volume was 500 mL at a pressure of 760 torr, when the volume is compressed to 450 mL, what is the pressure?

Plug in the values: (760 torr)(500 mL) = P2(450 mL) 760 torr x 500 mL/450 mL = P2 844 torr = P2 The pressure is 844 torr after compression. Charles' Law: The Temperature-Volume Law This law states that the volume of a given amount of gas held at constant pressure is directly proportional to the Kelvin temperature. V Same as before, a constant can be put in: V / T = C As the volume goes up, the temperature also goes up, and vice-versa. Also same as before, initial and final volumes and temperatures under constant pressure can be calculated. V1 / T1 = V2 / T2 = V3 / T3 etc. Gay-Lussac's Law: The Pressure Temperature Law This law states that the pressure of a given amount of gas held at constant volume is directly proportional to the Kelvin temperature. P Same as before, a constant can be put in: P / T = C As the pressure goes up, the temperature also goes up, and vice-versa.

Also same as before, initial and final volumes and temperatures under constant pressure can be calculated. P1 / T1 = P2 / T2 = P3 / T3 etc. Avogadro's Law: The Volume Amount Law Amedeo Avogadro (1776-1856) Gives the relationship between volume and amount when pressure and temperature are held constant. Remember amount is measured in moles. Also, since volume is one of the variables, that means the container holding the gas is flexible in some way and can expand or contract. If the amount of gas in a container is increased, the volume increases. If the amount of gas in a container is decreased, the volume decreases.

V FW = gRT / PV ; V = (0.1000 g)(0.0821 L-atm / mole-K)(295.3 K) / (0.0928 atm)(0.256 L) = 102 g / mole FW of CHF2 = 51.0 g / mole ; 102 / 51.0 = 2 ; C2H2F4 If the equation above is rearranged further, g / V = P x FW / RT = density you get an expression of the density of the gas as a function of T and FW. EXAMPLE 3: Compare the density of He and air (average FW = 28 g/mole) at 25.0 oC and 1.00 atm.

dHe = (4.003 g / mole)(1.00 atm) / (0.0821 L-atm / mole-K)(298 K) = 0.164 g / L dair = (28.0 g / mole)(1.00 atm) / (0.0821 L-atm / mole-K)(298 K) = 1.14 g / L EXAMPLE 4: Compare the density of air at 25.0 oC and air at 1807 oC and 1.00 atm.

dHe = (28.0 g / mole)(1.00 atm) / (0.0821 L-atm / mole-K)(298 K) = 1.14 g / L dair = (28.0 g / mole)(1.00 atm) / (0.0821 L-atm / mole-K)(2080 K) = 0.164 g / L Partial Pressures Dalton's Law of Partial Pressures states that the total pressure of a mixture of nonreacting gases is the sum of their individual partial pressures. Ptotal = Pa + Pb + Pc + ... Ptotal = naRT / V + nbRT / V + ncRT / V + ... or Ptotal = (na+ nb+ nc+ ... )RT / V The pressure in a flask containing a mixture of 1 mole of 0.20 mole O2 and 0.80 mole N2 would be the same as the same flask holding 1 mole of O2. Partial pressures are useful when gases are collected by bubbling through water (displacement).

The gas collected is saturated in water vapor which contibutes to the total number of moles of gas in the container. EXAMPLE 5: A sample of H2 was prepared in the laboratory by the reaction: Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g) 456 mL of gas was collected at 22.0 oC. The total pressure in the flask was 742 torr. How many moles of H2 were collected? The vapor pressure of H2O at 22.0 oC is 19.8 torr. nH2 = PH2V / RT ; nH2 = (0.9503 atm)(0.456 L) / (0.0821 L-atm / mole-K)(295 K) = 0.0179 mole H2. Non-Ideal Gases Johannes Diderik van der Waals (1837-1923) The ideal gas equation (PV=nRT) provides a valuable model of the relations between volume, pressure, temperature and number of particles in a gas.

As an ideal model it serves as a reference for the behavior of real gases. The ideal gas equation makes some simplifying assumptions which are obviously not quite true. Real molecules do have volume and do attract each other. All gases depart from ideal behavior under conditions of low temperature (when liquefaction begins) and high pressure (molecules are more crowed so the volume of the molecule becomes important). Refinements to the ideal gas equation can be made to correct for these deviations. In 1873 J. D. van der Waals proposed his equation, known as the van der Waals equation. As there are attractive forces between molecules, the pressure is lower than the ideal value. To account for this the pressure term is augmented by an attractive force term a/V2. Likewise real molecules have a volume. The volume of the molecules is represented by the term b. The term b is a function of a spherical diameter d known as the van der Waals diameter. The van der Waals equation for n moles of gas is: The values for a and b, below, are determined empirically: Charles's Law: Volume vs. Temperature of a Gas at Constant Pressure Charles's Law: Volume vs.

Temperature of a Gas at Constant Pressure Minor injury possible: this experiment involves heating water on a stove. Adult supervision recommended. Abstract This is a modern version of a classic experiment by Jacques Charles on the volume of a gas at different temperatures. Charles discovered the relationship between volume and temperature of gases that now bears his name. This project shows you a simple method for re-creating this famous experiment. Objective The goal of this project is to measure the relationship between the volume of a gas and its temperature, when the pressure of the gas is held constant. Credits Andrew Olson, Ph.D., Science Buddies Sources Gabel, Dorothy, 1996. "Learning Sequence Item 929: Gas Laws" in Scope, Sequence and Coordination: A National Curriculum Project for High School Science Education. Arlington, VA: National Science Teachers Association. Available online at: . Gardner, Robert, 1990. Famous Experiments You Can Do. New York, NY: Franklin Watts (pp. 103–104). Blauch, D., 2004. "Gas Laws: Pressure", Department of Chemistry, Davidson College. [accessed January 23, 2006] .

Grateful acknowledgment is also made to Prof. Blauch, Department of Chemistry, Davidson College for making his Gas Law and other Chemistry Applets freely available on the Web. Cite This Page MLA Style Science Buddies Staff. "Charles's Law: Volume vs. Temperature of a Gas at Constant Pressure" Science Buddies. Science Buddies, 3 Oct. 2014. Web. 19 Jan. 2017 < APA Style Science Buddies Staff. (2014, October 3). Charles's Law: Volume vs. Temperature of a Gas at Constant Pressure. Retrieved January 19, 2017 from Share your story with Science Buddies! I Did This Project! Please log in and let us know how things went. Last edit date: 2014-10-03 Introduction This is a modern version of a classic experiment by Jacques Charles (who was also interested in flying balloons). Charles studied the volume of a sample of air—sealed in a glass tube with a U-shaped curve—as he systematically changed the temperature by immersing the tube in a water bath. The air was trapped by a column of mercury, added to the open end of the tube.

By changing the amount of mercury in the tube, Charles could maintain a constant pressure on the trapped air as the temperature was changed. Charles's apparatus was an example of a manometer, a device used to measure pressure. The following diagram and description, from Prof. David N. Blauch, of Davidson College, explain how a manometer works. "A manometer is a device employed to measure pressure. There are a variety of manometer designs. A simple, common design is to seal a length of glass tubing and then bend the glass tube into a U-shape. The glass tube is then filled with a liquid, typically mercury, so that all trapped air is removed from the sealed end of the tube. The glass tube is then positioned with the curved region at the bottom. The mercury settles to the bottom (see the picture at the left). "After the mercury settles to the bottom of the manometer, a vacuum is produced in the sealed tube (the left tube in the picture). The open tube is connected to the system whose pressure is being measured. In the sealed tube, there is no gas to exert a force on the mercury. In the tube connected to the system, the gas in the system exerts a force on the mercury.

The net result is that the column of mercury in the left (sealed) tube is higher than that in the right (unsealed) tube. The difference in the heights of the columns of mercury is a measure of the pressure of gas in the system. "In the example at the left, the top of the left column of mercury corresponds to 875 mm on the scale. The top of the right column of mercury corresponds to 115 mm. The difference in heights is 875 mm − 115 mm = 760. mm, which indicates that the pressure is 760. mm Hg or 760. torr." (Blauch, 2004). You can learn more about how manometers work, and even run a simulated Charles's Law experiment by visiting the Chemistry Applet website (see Bibliography). This would be excellent preparation for doing the experiment on your own, so we highly recommend it. Note: to calculate the volume of gas in the applets, you will need to know that the inside diameter of the applet's manometer tube is 4.286 cm.

You can repeat Charles's experiments for yourself with an inexpensive, modern apparatus based on a disposable plastic syringe and a water bath. (Mercury is a dangerous neurotoxin, so we'll avoid working with it.) Terms and Concepts To do this project, you should do research that enables you to understand the following terms and concepts: pressure, cooking pot, deeper than syringe is tall stove Disclaimer: Science Buddies occasionally provides information (such as part numbers, supplier names, and supplier weblinks) to assist our users in locating specialty items for individual projects. The information is provided solely as a convenience to our users. We do our best to make sure that part numbers and descriptions are accurate when first listed. However, since part numbers do change as items are obsoleted or improved, please send us an email if you run across any parts that are no longer available. We also do our best to make sure that any listed supplier provides prompt, courteous service. Science Buddies does participate in affiliate programs with Amazon.com , Carolina Biological , Jameco Electronics , and AquaPhoenix Education . Proceeds from the affiliate programs help support Science Buddies, a 501(c)(3) public charity.

If you have any comments (positive or negative) related to purchases you've made for science fair projects from recommendations on our site, please let us know. Write to us at scibuddy@sciencebuddies.org . Share your story with Science Buddies! I Did This Project! Please log in and let us know how things went. Experimental Procedure Experimental Apparatus Diagram showing how to set up syringe. The thin wire between the plunger tip and the inner syringe wall allows air to escape from in front of the plunger in order to equalize pressure. It is removed before starting the experiment. Diagram adapted from Gabel, 1996. Before starting the experiment, do your background research so that you are knowledgeable about the terms, concepts and questions, above. With the plunger removed from the syringe, seal the tip of the syringe with a tight-fitting cap. If a suitable cap is not available, you can try epoxy or silicone sealant. Allow the epoxy or silicone the recommended curing time before proceeding with the experiment. (Note: if you seal the tip with the plunger in place, you will probably not be able to remove the plunger unless you destroy the seal.

Why?) When your sealed syringe is ready for use, insert the plunger to the 20 mL mark of the syringe along with a thin wire as shown in the diagram above. The wire will allow air to escape from beneath the plunger, equalizing the pressure in the syringe with the atmosphere. Use the lower ring of the plunger as your indicator. Hold the plunger in place and carefully withdraw the wire. Make sure that the plunger can move freely in the syringe, and that the tip of the syringe is well-sealed. Give the plunger a small downward push, and verify that it springs back. If it does not, you may need to lubricate the side of the plunger with a small amount of silicone lubricant or you may not have sealed the tip of your syringe properly. When you are satisfied with the results of the previous step, record the initial volume of air in the syringe and the ambient temperature. You will be immersing the syringe into a water bath, and observing the changes in volume of the gas as you change the temperature of the water. Since the air in the syringe will make it buoyant, you need a way to hold the syringe under the water.

If you have a ringstand and clamp, you're all set. Otherwise, you can put together a homemade clamp with materials you'll probably have around the house. Here's how: Wrap a rubber band around the top of the syringe tube, just below the finger flanges. Insert the chopsticks (as noted in Materials & Equipment, wood dowels can be substituted for chopsticks) through loops of this rubber band, one on either side of the syringe. Slide the syringe so that it is about 7–8 cm (3 in) in from the ends of the chopsticks. Wrap the second rubber band around the short ends of the chopsticks. This will make a "V" shape, with the syringe held tightly down near the point. This second rubber band can also be used to hold the thermometer upright in the water. Keep the bulb immersed in the water, but not touching the side or bottom of the pot. Place this assembly on the top of your cooking pot, so that the chopsticks are supported by the rim of the pot and the syringe sticks down into the pot.

To hold the syringe in place when the pot is filled with water, place your weight (e.g., a can of soup) on top of the wide end of the "V" made by the chopsticks. Make any necessary adjustments to make the syringe and thermometer stable, and make sure that you can read the scale on the syringe. Making the Measurements and Presenting Your Results Remove the syringe and thermometer assembly from the pot and set them aside. Place the pot on the stove, but don't turn on the burner yet. Fill the pot with ice cubes and enough water to immerse the syringe to somewhere between the 25 and 30 mL marks. Replace the syringe and thermometer assembly, and weight it down securely. Allow several minutes temperature in the water bath to stabilize and for the temperature of the air in the syringe to equilibrate with the water bath. Gentle stirring may help, but be careful not to break the thermometer or knock your weight off your clamp. Record the temperature of the water bath and the volume of the air in the syringe. You may want to tap the plunger lightly to make sure it is free to move.

(If necessary, carefully (and briefly) lift the syringe out of the water to read the volume. You may want to have an adult help you with this part.) Turn the burner on (no higher than medium heat) to gradually heat the water. At regular intervals (e.g., every 10°C), turn the heat off and allow the temperature to stabilize. Again, record the temperature of the water bath and the volume of air in the syringe. Repeat the previous step up to 80 or 90°C. The pot will be quite full, so it is best to avoid boiling the water. As with any experiment, it is a good idea to repeat your measurements to be sure that your results are consistent. We suggest at least three separate trials. (Note that the temperatures used do not need to be exactly the same from trial to trial!) Make a graph of gas volume vs. temperature for all of your data points. It's a good idea to use a different symbol for each of your trials (if something was wrong with one particular trial, it may help you understand what went wrong). Questions What is the relationship between volume and temperature in your data set?

Can you extrapolate from your data to find the temperature that corresponds to a gas volume of zero? How confident are you with this result, and why? Would your data look different if you used kelvins for the temperature axis instead of degrees Celsius? Was the assumption of constant pressure valid? What are the possible sources of error in your experiment? Share your story with Science Buddies! I Did This Project! Please log in and let us know how things went. Variations For a more advanced version of this project, combine it with the Boyle's Law project (see above) and do background research on statistical mechanics, and explain your results in terms of molecular motions. Share your story with Science Buddies! I Did This Project! Please log in and let us know how things went. Ask an Expert The Ask an Expert Forum is intended to be a place where students can go to find answers to science questions that they have been unable to find using other resources. If you have specific questions about your science fair project or science fair, our team of volunteer scientists can help. Our Experts won't do the work for you, but they will make suggestions, offer guidance, and help you troubleshoot. Boyle's Law: Pressure vs.

Volume of a Gas at Constant Temperature Boyle's Law: Pressure vs. Volume of a Gas at Constant Temperature Minor injury possible: be careful when stacking bricks on syringe apparatus. Abstract This is a modern version of a classic experiment by Robert Boyle on the compressibility of gases. Boyle discovered the relationship between pressure and volume of gases that now bears his name. This project shows you a simple method for re-creating this famous experiment. Objective The goal of this project is to measure the relationship between the volume of a gas and its pressure, when the temperature of the gas is held constant. Credits Written by: Andrew Olson, Ph.D., Science Buddies Sources Gabel, Dorothy, 1996. "Learning Sequence Item 929: Gas Laws" in Scope, Sequence and Coordination: A National Curriculum Project for High School Science Education. Arlington, VA: National Science Teachers Association. Available online at: . Gardner, Robert, 1990. Famous Experiments You Can Do. New York, NY: Franklin Watts (pp. 61–65). Blauch, D., 2004. "Gas Laws: Pressure", Department of Chemistry, Davidson College.

[accessed December 13, 2005] . Grateful acknowledgment is also made to Prof. Blauch, Department of Chemistry, Davidson College for making his Gas Law and other Chemistry Applets freely available on the Web. Cite This Page MLA Style Science Buddies Staff. "Boyle's Law: Pressure vs. Volume of a Gas at Constant Temperature" Science Buddies. Science Buddies, 3 Oct. 2014. Web. 19 Jan. 2017 < APA Style Science Buddies Staff. (2014, October 3). Boyle's Law: Pressure vs. Volume of a Gas at Constant Temperature. Retrieved January 19, 2017 from Share your story with Science Buddies! I Did This Project! Please log in and let us know how things went. Last edit date: 2014-10-03 Introduction This project re-creates a study begun in 1662 by Robert Boyle. Now that's a classic! Unlike liquids, gases are compressible. Boyle systematically studied the compression of air, sealed in a glass tube with a U-shaped curve.

The air was trapped by a column of mercury, added to the open end of the tube. By changing the amount of mercury in the tube, Boyle could change the pressure exerted on the trapped air. Boyle's apparatus was an example of a manometer, a device used to measure pressure. The following diagram and description, from Prof. David N. Blauch, of Davidson College, explain how a manometer works. "A manometer is a device employed to measure pressure. There are a variety of manometer designs. A simple, common design is to seal a length of glass tubing and then bend the glass tube into a U-shape. The glass tube is then filled with a liquid, typically mercury, so that all trapped air is removed from the sealed end of the tube. The glass tube is then positioned with the curved region at the bottom. The mercury settles to the bottom (see the picture at the left). "After the mercury settles to the bottom of the manometer, a vacuum is produced in the sealed tube (the left tube in the picture). The open tube is connected to the system whose pressure is being measured. In the sealed tube, there is no gas to exert a force on the mercury.

In the tube connected to the system, the gas in the system exerts a force on the mercury. The net result is that the column of mercury in the left (sealed) tube is higher than that in the right (unsealed) tube. The difference in the heights of the columns of mercury is a measure of the pressure of gas in the system. "In the example at the left, the top of the left column of mercury corresponds to 875 mm on the scale. The top of the right column of mercury corresponds to 115 mm. The difference in heights is 875 mm − 115 mm = 760. mm, which indicates that the pressure is 760. mm Hg or 760. torr." (Blauch, 2004). You can learn more about how manometers work, and even run a simulated Boyle's Law experiment by visiting the Chemistry Applet website (see Bibliography). This would be excellent preparation for doing the experiment on your own, so we highly recommend it. You can repeat Boyle's experiments with an inexpensive, modern apparatus based on a disposable plastic syringe.

Mercury is a dangerous neurotoxin, so we'll avoid working with it. Instead, you will compress the air in the syringe with . . . bricks! Pressure is force exerted over a unit area, so the units are those of force divided by area. The larger area of a syringe (compared to a narrow glass tube) means that you will need a larger force to compress the gas. Bricks, which are stackable and weigh about 1 kg apiece, will do the job nicely. Terms and Concepts To do this project, you should do research that enables you to understand the following terms and concepts: pressure, Questions What assumption is made about the temperature of the gas in this experiment? Pressure is measured in many different units. For example, blood pressure is typically expressed in mm of Hg (mercury), as in Boyle's experiments. In the U.S., pressure for car tires is typically expressed in pounds per square inch. The SI unit for pressure is the pascal (Pa). How are all of these different units related? How can you convert from one to another? Bibliography Here is the sequence of three Chemistry Applet webpages mentioned in the Introduction.

These will really help your understanding of Boyle's Law if you take the time to do the virtual experiments! Blauch, D., 2004. "Gas Laws: Pressure", Department of Chemistry, Davidson College. [accessed December 13, 2005] , Blauch, D., 2004. "Gas Laws: Boyle's Law", Department of Chemistry, Davidson College. [accessed December 13, 2005] , and Blauch, D., 2004. "Gas Laws: Boyle's Law Calculations", Department of Chemistry, Davidson College. [accessed December 13, 2005] . The Sizes.com website has an exhaustive index of units of measure, including the SI definition for the pascal, and a link to a calculator to convert between various units of pressure: Editor, Sizes.com, 2000. "pascal" [accessed December 13, 2005] News Feed on This Topic , , Note: A computerized matching algorithm suggests the above articles. It's not as smart as you are, and it may occasionally give humorous, ridiculous, or even annoying results!

Learn more about the News Feed Materials and Equipment Syringe that can hold at least 30 mL. If the syringe does not come with a tight-fitting cap, you will also need epoxy or silicone sealant. A suitable syringe that comes with a cap is available from Amazon.com . Wood blocks with center holes drilled partway through (2); see the diagram in the Experimental Procedure section. The first block will hold the syringe upright, and will need a hole that is just slightly larger than the diameter of the syringe, plus a smaller hole to accommodate the syringe tip. For greater stability, this block can be clamped in place (optional). The second block will be placed on top of the syringe plunger, and will act as a shelf for the bricks; the diameter of the hole should fit the top of the syringe plunger snugly. A small piece of wire Bricks with approximately uniform size and weight of about 1 kg (4 or 5 bricks) Scale for weighing bricks; a bathroom scale should be adequate. Graph paper Calculator Optional: "C" clamps (2). These are available at hardware stores and online through suppliers such as Amazon.com .

Lab notebook Disclaimer: Science Buddies occasionally provides information (such as part numbers, supplier names, and supplier weblinks) to assist our users in locating specialty items for individual projects. The information is provided solely as a convenience to our users. We do our best to make sure that part numbers and descriptions are accurate when first listed. However, since part numbers do change as items are obsoleted or improved, please send us an email if you run across any parts that are no longer available. We also do our best to make sure that any listed supplier provides prompt, courteous service. Science Buddies does participate in affiliate programs with Amazon.com , Carolina Biological , Jameco Electronics , and AquaPhoenix Education . Proceeds from the affiliate programs help support Science Buddies, a 501(c)(3) public charity. If you have any comments (positive or negative) related to purchases you've made for science fair projects from recommendations on our site, please let us know. Write to us at scibuddy@sciencebuddies.org . Share your story with Science Buddies! I Did This Project! Please log in and let us know how things went.

Experimental Procedure Making the Measurements Diagram of experimental setup showing syringe with sealed tip held vertically in pre-drilled wood block support. The thin wire between the plunger tip and the inner syringe wall allows air to escape from in front of the plunger in order to equalize pressure. It is removed before starting the experiment. The second pre-drilled wood block is placed atop the syringe plunger and acts as a shelf for bricks to increase the pressure on the plunger. Diagram from Gabel, 1996. Before starting the experiment, do your background research so that you are knowledgeable about the terms, concepts and questions, above. With the plunger removed from the syringe, seal the tip of the syringe with a tight-fitting cap. If a suitable cap is not available, you can try epoxy or silicone sealant. Allow the epoxy or silicone the recommended curing time before proceeding with the experiment. (Note: if you seal the tip with the plunger in place, you will probably not be able to remove the plunger unless you destroy the seal.

Why?) When your sealed syringe is ready for use, insert the syringe firmly, tip down, into the pre-drilled hole in the bottom wood block support, as shown in the diagram. The syringe should fit snugly, so it does not wobble when you load it up with bricks. You may wish to clamp the block in place. (Note: clamp to a workbench, not a piece of fine furniture!) Insert the plunger to the 30 ml mark of the syringe along with a thin wire as shown in the diagram. The wire will allow air to escape from beneath the plunger, equalizing the pressure in the syringe with the atmosphere. Use the lower ring of the plunger as your indicator. Hold the plunger in place and carefully withdraw the wire. Make sure that the plunger can move freely in the syringe, and that the tip of the syringe is well-sealed. Give the plunger a small downward push, and verify that it springs back. If it does not, you may need to lubricate the side of the plunger with a small amount of silicone lubricant or you may not have sealed the tip of your syringe properly.

When you are satisfied with the results of the previous step, record the initial volume of air in the syringe. Place second wood block over the top of the plunger, as shown in the diagram. This wood block will act as a shelf to hold bricks in order to exert downward force on the plunger. Make sure that the shelf is level and well-seated on the plunger. Place the first brick on the shelf. You may need to tap on the brick to free the plunger. Note the resulting volume of air in the syringe. Repeat the previous step until you have 4 or 5 bricks stacked on the syringe. With each added brick, note the volume of air, and the number of bricks. Next, you will remove the bricks, one at a time, noting the volume of air in the syringe each time. Again, you may need to tap on the shelf to free the plunger. Take the average of your two values (from loading and unloading) for each number of bricks. Remove the plunger and repeat steps 4–12 so that you have at least 5 trials.

Analyzing Your Data and Converting to Standard Units Calculate the average and standard deviation of the volume for each of your data points over the repeated trials. At this point, you have measured the volume of air in the syringe as a function of the number of bricks pushing down on the plunger. The next step is to convert from bricks to units of pressure. Pressure is defined as force per unit area. The SI unit for pressure is the pascal (Pa), which is defined as the force of 1 newton acting over an area of 1 square meter. So you will need to know the force exerted by the bricks, plus the area of the plunger. You can calculate the downward force, F of the brick(s) by multiplying the mass, m, of the brick(s) times the acceleration, a, due to gravity (F = ma), where a = 9.83 N/kg. You can calculate the area of the plunger (in units of square meters) by measuring the diameter, and recalling the formula for calculating the area of a circle. You can then calculate the pressure, in Pa, by dividing the force, F, generated by each stack of bricks by the area, A, of the plunger.

Presenting Your Results Boyle found that, when temperature is held constant, the pressure and volume of a gas are inversely related. In mathematical terms, we can write PV = k, where P, naturally, is pressure, and V is volume. Alternatively, we can write P1V1 = P2V2, where the subscripts indicate measurements of different pressure-volume pairs. The equation shows that as pressure increases, volume must decrease, and vice versa. Do your results agree? What method(s) of graphing your results will show the inverse relationship between pressure and volume? (If you need ideas, check out the virtual experiment links in the Bibliography section.) You should be able to describe in your own words how your graph illustrates an equation for Boyle's Law. What is the pressure on the air in the syringe before you start adding weight to the the plunger? Can you extrapolate from your measurements to give a value for this pressure? Can you think of a way to check the results of your extrapolation? Questions For a more advanced version of this project, combine it with the Charles's Law project (see above) and do background research on statistical mechanics, and explain your results in terms of molecular motions. Share your story with Science Buddies!

I Did This Project! Please log in and let us know how things went. Ask an Expert The Ask an Expert Forum is intended to be a place where students can go to find answers to science questions that they have been unable to find using other resources. If you have specific questions about your science fair project or science fair, our team of volunteer scientists can help. Our Experts won't do the work for you, but they will make suggestions, offer guidance, and help you troubleshoot.

Question: Which scientific law describes the relationship between the volume and temperature of a gas at constant pressure?
Answer: Charles's Law